Organic Chemistry Lecture 1
Lewis structures, bonds, hybridization and (lewis) acids and bases
Reference: Bruice, P.Y., Organic Chemistry, 8th edition; Pearson Education: Harlow,
Essex, UK
Chapter 1
- Organic compounds = compounds that are based on carbon
- Inorganic compounds = compounds derived from minerals and have no vital force.
Why organic chemistry?
-> All of the compounds that make life possible are organic compounds.
1.1 The Structure of an Atom
The nucleus contains the protons and neutrons (and is thus positively charged); the electrons are in
the electron cloud (negatively charged). The amount of protons and electrons must be the same to
get an uncharged atom.
Electrons are always moving and therefore have a kinetic energy. This kinetic energy counteracts
the attractive force of the protons in the nucleus (otherwise the electrons would be pulled towards
the nucleus).
Protons and neutrons have about the same mass and are 1800 times more massive than
electrons. Most of the mass of an atom is in the nucleus, however the volume is mostly in the
electron cloud.
The atomic number is the number of protons in the nucleus and therefore also the amount of
electrons in an uncharged atom. The mass number is the sum of the protons and neutrons. There are
isotopes present who have the same amount of protons but not the same amount of neutrons.
Atomic mass is the weighted average of the isotopes of an element and the molecular mass is the
sum of the atomic masses in a molecule.
1.2 Covalent Bonds
According to Lewis’s theory, an atom will share, donate or accept electrons to achieve a full shell, this
is called the octet rule.
Achieving a Filled Outer Shell by Losing or Gaining Electrons:
Those with 1 valence electron can lose an electron to achieve a full shell (under laying shell). They
become positively charged.
Those with 7 valence electrons can accept an electron to achieve a full shell. They become negatively
charged.
Hydrogen atoms can lose an electron to become a proton (H +) or accept an electron to become a
hydride ion (H-).
Achieving a Filled Outer Shell by Sharing Electrons:
Atoms can share electrons to fulfill the octet rule.
Nonpolar and Polar Covalent Bonds:
if the EN difference is smaller than 0.5, the bond is nonpolar. If the EN difference is between 0.5 and
1.9, the bond is polar. The direction of the bond polarity goes from the δ+ (the smallest EN) to the δ-
(the biggest EN). If the EN difference is bigger than 1.9, the electron is held together by electrostatic
attraction rather than covalent bonds.
1
,Dipole Moments of Bonds:
A dipole moment is present in polar covalent bonds and is equal to the magnitude of the charge on
either atom: dipole moment = size of the charge * distance between the charges. It’s unit is debye
(D).
Electrostatic Potential Maps:
Electrostatic potential maps show how the charge is distributed in the molecule. Red indicates the
most negative electrostatic potential and blue indicates the most positive electrostatic potential.
1.3 How The Structure of a Compound is Represented
Lewis structure
- Lone pairs: nonbinding valence electrons
- Formal charge: the difference between the number of valence electrons on an atom when
it’s bonded and it’s own number of valence electrons. It’s own number of valence electrons =
half of the shared bonds and all lone pairs on the atom.
Formal charge = number valence electrons – (number of lone pairs + number bonds)
2
,1.4 Atomic Orbitals
An atomic orbital is a 3D region where the electron is most-likely found. According to the Heisenberg
uncertainty principle, it is impossible to determine the precise location and moment of an atomic
particle. The orbitals describe it’s probable location.
s Atomic Orbitals
p Atomic Orbitals
- Nodal plane = the node that passes through the center of the nucleus between the lobes
2.6 An Introduction to Molecular Orbital Theory (see SAM)
Molecular Orbital (MO) theory combines the tendency of atoms to form covalent bonds and their
wave-like properties. According to this theory, covalent bonds result when the atomic orbitales form
molecular orbitals.
3
, Forming a Sigma (σ) Bond and a pi (π) Bond
Overlapping s orbitals form a sigma bond.
Overlapping p orbitals form a pi bond.
Bonding and Antibonding Molecular Orbitals
4
Lewis structures, bonds, hybridization and (lewis) acids and bases
Reference: Bruice, P.Y., Organic Chemistry, 8th edition; Pearson Education: Harlow,
Essex, UK
Chapter 1
- Organic compounds = compounds that are based on carbon
- Inorganic compounds = compounds derived from minerals and have no vital force.
Why organic chemistry?
-> All of the compounds that make life possible are organic compounds.
1.1 The Structure of an Atom
The nucleus contains the protons and neutrons (and is thus positively charged); the electrons are in
the electron cloud (negatively charged). The amount of protons and electrons must be the same to
get an uncharged atom.
Electrons are always moving and therefore have a kinetic energy. This kinetic energy counteracts
the attractive force of the protons in the nucleus (otherwise the electrons would be pulled towards
the nucleus).
Protons and neutrons have about the same mass and are 1800 times more massive than
electrons. Most of the mass of an atom is in the nucleus, however the volume is mostly in the
electron cloud.
The atomic number is the number of protons in the nucleus and therefore also the amount of
electrons in an uncharged atom. The mass number is the sum of the protons and neutrons. There are
isotopes present who have the same amount of protons but not the same amount of neutrons.
Atomic mass is the weighted average of the isotopes of an element and the molecular mass is the
sum of the atomic masses in a molecule.
1.2 Covalent Bonds
According to Lewis’s theory, an atom will share, donate or accept electrons to achieve a full shell, this
is called the octet rule.
Achieving a Filled Outer Shell by Losing or Gaining Electrons:
Those with 1 valence electron can lose an electron to achieve a full shell (under laying shell). They
become positively charged.
Those with 7 valence electrons can accept an electron to achieve a full shell. They become negatively
charged.
Hydrogen atoms can lose an electron to become a proton (H +) or accept an electron to become a
hydride ion (H-).
Achieving a Filled Outer Shell by Sharing Electrons:
Atoms can share electrons to fulfill the octet rule.
Nonpolar and Polar Covalent Bonds:
if the EN difference is smaller than 0.5, the bond is nonpolar. If the EN difference is between 0.5 and
1.9, the bond is polar. The direction of the bond polarity goes from the δ+ (the smallest EN) to the δ-
(the biggest EN). If the EN difference is bigger than 1.9, the electron is held together by electrostatic
attraction rather than covalent bonds.
1
,Dipole Moments of Bonds:
A dipole moment is present in polar covalent bonds and is equal to the magnitude of the charge on
either atom: dipole moment = size of the charge * distance between the charges. It’s unit is debye
(D).
Electrostatic Potential Maps:
Electrostatic potential maps show how the charge is distributed in the molecule. Red indicates the
most negative electrostatic potential and blue indicates the most positive electrostatic potential.
1.3 How The Structure of a Compound is Represented
Lewis structure
- Lone pairs: nonbinding valence electrons
- Formal charge: the difference between the number of valence electrons on an atom when
it’s bonded and it’s own number of valence electrons. It’s own number of valence electrons =
half of the shared bonds and all lone pairs on the atom.
Formal charge = number valence electrons – (number of lone pairs + number bonds)
2
,1.4 Atomic Orbitals
An atomic orbital is a 3D region where the electron is most-likely found. According to the Heisenberg
uncertainty principle, it is impossible to determine the precise location and moment of an atomic
particle. The orbitals describe it’s probable location.
s Atomic Orbitals
p Atomic Orbitals
- Nodal plane = the node that passes through the center of the nucleus between the lobes
2.6 An Introduction to Molecular Orbital Theory (see SAM)
Molecular Orbital (MO) theory combines the tendency of atoms to form covalent bonds and their
wave-like properties. According to this theory, covalent bonds result when the atomic orbitales form
molecular orbitals.
3
, Forming a Sigma (σ) Bond and a pi (π) Bond
Overlapping s orbitals form a sigma bond.
Overlapping p orbitals form a pi bond.
Bonding and Antibonding Molecular Orbitals
4
